1. Chapter 6-Chemical Bonding
6.1-Introduction to Chemical Bonding
6.2-Covalent Bonding & Molecular Compounds
6.3-Ionic Bonding & Ionic Compounds
6.4-Metallic Bonding
6.5-Molecular Geometry

2. Warm-Up 12/1/10
1. Chlorine has an electronegativity value of 3.0. What type of bond will form between two Cl atoms?
2. Sodium has an electronegativity value of 0.9 and Chlorine has a value of 3.0. What type of bond will form between the two atoms?
3. How many valence electrons are in the outer shell of F, Be, and C?

3. Warm-Up 12/2/10 1. Write the electron dot notation for: a. Bi c. He
b. Sr d.Te
2. What are the exceptions to the octet rule?

4. Warm-Up 12/4/08 – Write questions
1. What gas was produced during the lab?
2. Why did the acid remain on the bottom of gas collection tubes when adding water?
3. What happened to the magnesium that reacted with hydrochloric acid?
4. What law explains the reason why the magnesium wasn’t destroyed during the reaction.
5. If you spill acid on your skin, what should you immediately do?

5. Warm-Up 1. What type of bond will form between: Na and Co Li and N
Ca and H
2. How many valence electrons do the following have:
a. Zn
b. B

6. Warm-Up 11/29
1.Define Electronegativity. 2.Which of the following is more electronegative: Mg and S C and O 3.What is a compound? When finished, turn in warm-up and pick up a paper from the side lab table. You will have 15 minutes to finish this paper.

7. 6.1-Introduction to Chemical Bonding
Pages

8. Why do bonds form?
Atoms are lazy! They want to have the lowest possible potential energy.
As independent particles, atoms have relatively high potential energy.
By bonding, atoms lower their potential energy and become more stable.

9. Types of Chemical Bonding
Chemical bond-a mutual electrical attraction between the nuclei & valence electrons of different atoms that binds the atoms together.
Major types of bonding:
Covalent bonding-sharing of electron pairs
Ionic bonding-electrical attraction between cations and anions
Metallic bonding-sharing of electrons in an “electron sea”

10. What type of bond will form?
Use electronegativity values to determine
Remember that electronegativity is a measure of an atom’s ability to attract electrons.
Bond Type
EN Difference
% Ionic
Nonpolar covalent
0-0.3
0-5
Polar covalent
5-50
Ionic
>1.7
50-100
Figure 6-2, pg. 162

11. What type of bond will form?
Nonpolar covalent-a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
Polar covalent-a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons
Ionic bond-a bond in which electron(s) are completely transferred from one atom to another, forming ions.

12. Comparing Electron Density in a Nonpolar & Polar Covalent Bond

13. What type of bond will form? examples
Bonding between sulfur and
EN difference
Bond type
More-negative atom
Hydrogen
0.4
Polar covalent
Sulfur
Cesium
1.8
Ionic
Chlorine
0.5

14. What type of bond will form? examples
Bonding between chlorine and
EN difference
Bond type
More-negative atom
Calcium
2.0
Ionic
Chlorine
Oxygen
0.5
Polar covalent
Bromine
0.2
Nonpolar covalent

15. 6.2-Covalent Bonding & Molecular Compounds
Pages

16. Vocab
Molecule-a neutral group of atoms that are held together by covalent bonds
Molecular compound-a compound whose simplest units are molecules
Example-water
Molecular formula-shows the types and numbers of atoms combined in a single molecule
Example-H2O
Diatomic molecule-a molecule containing only two atoms

17. Energy of Bond Formation
Potential Energy
based on position of an object
low PE = high stability

18. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
no interaction
increased attraction

19. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
increased repulsion
balanced attraction & repulsion

20. Energy of Bond Formation
Bond energy-energy required to break a bond
Bond length-distance between two bonded atoms at their minimum PE
Bond Energy
Bond Length

21. Energy of Bond Formation
Bond Energy
Short bond = high bond energy

22. Octet Rule All atoms want to be a noble gas. Why?
Noble gases have a full outer energy level. This gives great stability!
When atoms bond, they gain stability by filling their outer energy level through sharing or exchanging electrons.
Octet rule-chemical compounds tend to form so that each atom, but gaining, losing, or sharing electrons, has an octet of electrons in it highest occupied energy level.

23. Bonding electron pair in overlapping orbitals
Orbitals Overlap 1s 2s
2px
2py
2pz F
Fluorine atoms 1s 2s
2px
2py
2pz F
Fluorine molecule
Bonding electron pair in overlapping orbitals

24. Exceptions to the Octet Rule
Hydrogen (H)-2
Beryllium (Be)-4
Boron (B)-6
Some atoms want more than 8 electrons. This is called expanded valence, and it involves bonding in the d orbitals.

25. Orbitals Overlap-Exception
Bonding electron pair in overlapping orbitals 1s H 2s
2px
2py
2pz Cl
Hydrogen chloride molecule 1s H 2s
2px
2py
2pz Cl
Hydrogen & Chlorine atoms

26. Electron Dot Notation
An electron configuration notation in which only the valence electrons of an atom are shown, indicated by dots around the element’s symbol.

27. Steps to Electron Dot Notation
Determine the # of valence electrons in an atom of the element.
Write the symbol for the element.
Place dots around the symbol’s sides to indicate the valence electrons.
Dots should be placed as far from each other as possible!

28. Electron Dot Notation Example 1
Draw the electron dot notation for an atom of oxygen.
# of valence electrons?
Symbol? 6 O O

29. Electron Dot Notation Example 2
Draw the electron dot notation for an atom of boron.
# of valence electrons?
Symbol? 3 B B

30. F F F Lewis Structures Used to represent molecules.
Combine electron dot notations to show sharing of electrons. F F F
Fluorine atoms
Fluorine molecule

31. Parts of a Lewis Structure
Unshared or lone pair-electrons not involved in bonding that belong only to one atom
A line represents two electrons that are shared, bonding the 2 atoms together

32. NAS Method to drawing Lewis Structures
Needed, Available, Shared Method
Determine how many electrons are available by adding up the total number of valence electrons each atom in the molecule has.
Determine how many electrons are needed to satisfy each atoms’ octet. Don’t forget exceptions!
Subtract available from needed. This will tell you how many electrons have to be shared.
Divide the shared electrons by 2 to figure out how many bonds you will need.

33. Don’t forget the lone pairs!
Always count how many e- you used &make sure it’s not more than you had available!
Draw the Lewis structure for water, H2O.
How many electrons are available (A)?
How many electrons are needed (N)?
How many electrons must be shared (S)?
How many bonds will this molecule have?
Draw it!
H-2x1=2
O-1x6=6
Total=8
H-2x2=4
O-1x8=8
Total=12
Bond=2 e- N A S 12 H O H 8 4
2=2 bonds
Don’t forget the lone pairs!

34. I H C H H NAS Method Example 2 Draw the Lewis structure for CH3I. N A
How many electrons are available (A)?
How many electrons are needed (N)?
How many electrons must be shared (S)?
How many bonds will this molecule have?
Draw it! H C H
C-1x4=4
H-3x1=3
I-1x7=7
Total=14
C-1x8=8
H-3x2=6
I-1x8=8
Total=22 H N A S 22 14 8
2=4 bonds

35. Helpful Hints
ALWAYS count the electrons that you used. You can’t use more or less than you had available!
If you have a carbon (C) atom, this will always be in the center. Otherwise, usually the least electronegative element will be the central atom.
Hydrogen is NEVER the central atom since it will only bond two atoms!

36. Multiple Bonds
A single line connecting 2 atoms represents 2 electrons being shared and is called a single bond.
Atoms can share 2, 4, or 6 electrons.
2=single bond
4=double bond
6=triple bond
X-X
X=X
X≡X

37. Multiple Bond Example 1 Draw the Lewis structure for ethene, CH2O.
Use NAS to determine how many bonds the molecule will have.
What atom will be the central atom?
Draw it!

38. O H C H Multiple Bond Example 1
Carbon in center, with other atoms arrange around it.
Single bond atoms together. Have you used all of your bonds?
Where can you put the last bond?
Remember octet exceptions! O N A S 20 H C H 12 8
2=4 bonds

39. N N Multiple Bond Example 2 Draw the Lewis structure for nitrogen, N2.16 N N 10 6
2=3 bonds

40. [ ] O S Resonance Structures Lewis structures for SO2
Which one is correct?
Neither is exactly right. True structure is actually “in between” the two.
This is called resonance, and is indicated by placing brackets around the structures and a double headed arrow between them.
Resonance only occurs when there are multiple bonds! [ ] O S

41. 6.3-Ionic Bonding & Ionic Compounds
Pages

42. Vocab
Ionic compound-composed of positive & negative ions that are combined so that the numbers of positive & negative charges are equal.
Formula unit-the simplest collection of atoms from which an ionic compound’s formula can be established.
The smallest ratio of atoms in an ionic compound

43. Formation of Ionic Compounds
Electrons are transferred in an ionic compound to create anions and cations which are attracted to each other. Na Cl + - Na Cl

44. Ionic Bonding All atoms bond to minimize their potential energy.
Ions form a structure called a crystal lattice when bonding.
This arrangement minimizes repulsive forces and maximizes attractive forces between ions.

45. Energy of Bond Formation
Lattice Energy
Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

46. Comparison of Molecular & Ionic Compounds
Forces holding ions together are much stronger than attractive forces between molecules. This accounts for:
Higher melting and boiling points-it take more energy to break the bonds holding the ions together
Hard but brittle-if the crystal lattice structure slips, repulsive forces “line up,” causing a break in the structure
Ionic compounds are electrical conductors when molten or dissolved. The ions can move freely about and carry an electric current.

47. Polyatomic ions
Polyatomic ion-a charged group of covalently bonded atoms
Has both molecular & ionic characteristics
Charge is due to a shortage (+) or excess (-) of electrons
When drawing Lewis structures, adjust the available (A) electrons, place the structure in brackets, and superscript the charge.

48. Lewis structure: Polyatomics
Draw the Lewis structure of the ammonium ion, NH4+. [ ] + H N A S 16 N 8 8
2=4 bonds

49. Lewis structure: Polyatomics
Draw the Lewis structure of the nitrate ion, NO3-.
This would also have 2 resonance structures. [ ] - O N A S 32 N 24 8
2=4 bonds

50. 6.4-Metallic Bonding
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51. Metallic bonding
In a metal, vacant outer orbitals overlap. Outer electrons can move freely through these overlapping orbitals.
These electrons are called delocalized, and form a “sea of electrons.”
Metallic bonding-chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

52. Metallic bonding
“Sea of electrons” accounts for most properties of metals:
Electrical conductivity
Thermal conductivity
Shiny appearance
Due to absorption of many light frequencies
Malleability-ability of a substance to be hammered or beaten into then sheets.
Ductility-ability of a substance to be drawn, pulled, or extruded through a small opening to produce a thin wire.

53. 6.5-Molecular Geometry
Pages

54. VSEPR Theory “Valence Shell Electron Pair Repulsion” Theory
Electron pairs orient themselves in order to minimize repulsive forces.

55. VSEPR Theory Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-
Molecular geometry-shape the molecule takes on to minimize the repulsion between its electrons.

56. VSEPR Theory Lone pairs reduce the bond angle between atoms.

57. Steps to Determine Molecular Shape
Draw the Lewis Diagram.
Tally up e- pairs & atoms bonded to central atom.
Shape is determined by the # of atoms and e- pairs bonded to central atom.
Know the 6 common shapes
& their bond angles!

58. Determining Molecular Shape
A=central atom
B=any atoms bonded to central atom
E=e- pairs belonging to central atom

59. Common Molecular Shapes: AB2
2 atoms bonded to central atom
No lone pairs
LINEAR
180°
BeH2

60. Common Molecular Shapes: AB2E
2 atoms bonded to central atom
1 lone pair
SO2
BENT
<120°

61. Common Molecular Shapes: AB3
3 atoms bonded to central atom
No lone pair
BF3
TRIGONAL PLANAR
120°

62. Common Molecular Shapes: AB4
4 atoms bonded to central atom
No lone pairs
CH4
TETRAHEDRAL
109.5°

63. Common Molecular Shapes: AB3E
3 atoms bonded to central atom
1 lone pair
NH3
TRIGONAL PYRAMIDAL
107°

64. Common Molecular Shapes: AB2E2
2 atoms bonded to central atom
2 lone pairs
H2O
BENT
104.5°

65. Steps to determining molecular geometry
Draw Lewis structure
Determine number of atoms and lone pairs bonded to central atom
Assign a shape based on step 2.

66. Molecular Geometry Shape Atoms bonded to central atom Lone pairs
Type of molecule
Linear 2
AB2
Bent/angular 1
AB2E
Trigonal-planar 3
AB3
Tetrahedral 4
AB4
Trigonal-pyramidal
AB3E
AB2E2

67. F P F F 107° Examples TRIGONAL PYRAMIDAL PF3
Atoms bonded to central atom? 3
Lone pairs? 1
AB3E
F P F F
PF3
TRIGONAL PYRAMIDAL
107°

68. O C O 180° Examples LINEAR CO2 Atoms bonded to central atom?
Lone pairs?
AB2
LINEAR
180°

69. Warm-Up 12/10/08 Predict the shapes of the following molecules: CH4
NH2Cl
Tetrahedral
Bent
Trigonal-pyramidal

70. + - Dipole Moment H Cl Direction of the polar bond in a molecule
Arrow points toward the more e-neg atom
H Cl + -

71. Determining Molecular Polarity
Depends on:
dipole moments
molecular shape

72. Determining Molecular Polarity
Nonpolar Molecules
Dipole moments are symmetrical and cancel out.
BF3 F B

73. Determining Molecular Polarity
Polar Molecules
Dipole moments are asymmetrical and don’t cancel .
H2O H O
net
dipole
moment

74. Determining Molecular Polarity
Therefore, polar molecules have...
asymmetrical shape (lone pairs) or
asymmetrical atoms
CHCl3 H Cl
net
dipole
moment

75. Intermolecular Forces
Intermolecular forces-forces of attraction between molecules
Vary in strength, but generally weaker than bonds
3 types:
Dipole-dipole forces
Hydrogen bonding
London Dispersion forces

76. Br-F F-F Dipole-dipole forces
Forces of attraction between 2 polar molecules
Short-range, act only between nearby molecules
Compound
Boiling point
Bromine fluoride
-20°C
Fluorine
-188°C
Br-F
F-F

77. Hydrogen bonding Very strong type of dipole-dipole force
Found in compounds that contain H-F, H-O, or H-N
Strong electronegativity difference makes bonds highly polar
Partially positive hydrogen is attracted to an unshared pair of electrons on an adjacent molecule.
Hydrogen bonding-intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

78. Hydrogen Bonding in Water

79. London Dispersion Forces
AKA “van der Waals” forces
LDF-the intermolecular attraction resulting from the constant motion of electrons and the creation of instantaneous dipoles
Electrons are always moving around, which sometimes causes uneven electron distribution. This creates temporary “instantaneous” dipoles.